CHEMISTRY TOPIC 10

10.1 Proton Number and the Periodic Table

Atoms:
- Smallest part of an element.
- Cannot be split into anything simpler.
Every element has its own type of atom.
- Example: All zinc atoms are alike, but zinc atoms are different from copper, gold, and iron atoms.
Subatomic particles of an atom:
- Protons, neutrons, and electrons.
- Protons and neutrons make up the nucleus (center of atom).
- Electrons whizz around the outside.
Structure of the Carbon Atom:
- 6 protons (+ charge)
- 6 neutrons (no charge)
- Electrons orbit the nucleus
Proton number = number of protons = atom number
Number of protons = number of electrons in a neutral atom
All elements have a different number of protons.
Characteristics of subatomic particles:
- Proton: Relative mass = 1, Charge = +1
- Neutron: Relative mass = 1, Charge = 0
- Electron: Relative mass = $1/1840$ , Charge = -1
Atoms are neutral because the number of protons equals the number of electrons.
The whole mass of the atom is concentrated in the nucleus because electrons are very light compared to protons and neutrons.
Classification of the elements in the Periodic Table.
- Arranged in ascending order of their proton numbers.

10.2 Electrons in Atoms

Electrons orbit the nucleus of an atom.
Electrons occupy shells/energy levels/orbits.
- First shell: maximum 2 electrons
- Second shell: maximum 8 electrons
- Third shell: maximum 8 electrons (in this case, the last, so-called Valence Shell.)
Rules to write electron configuration:
- The first shell (close to the nucleus) can hold 2 electrons.
- The second shell onward can hold 8 electrons.
- Electrons fill the first shell first; the remaining electrons will go to the next shell.
Electron configuration = Electron arrangement.
Electron arrangement gives them their chemical properties.
- Outermost electrons determine where it can and cannot form bonds with other atoms.
- Determines its stability and what kind of reactions it will undergo.
Elements from the same group have the same amount of electrons at their outermost/valence shell.
Elements from the same period have the same number of shells.

10.3 Making Ions

Group 8 elements:
- Stable and chemically unreactive.
- Known as inert gases.
- Example: Helium (He) and Neon (Ne)
All the elements in group 8 have eight electrons (two for helium) in their outermost/valence shell.
It is the stable electron arrangement / stable structure.
It will maintain the electron arrangements, and they will not gain, lose, or share electrons with other atoms.
Group 8 elements will not combine with each other to form molecules or with other atoms to form compounds.
- Example: $Ne2$ , $He2$ , $O2$ , $Cl2$
Besides Group 8, atoms from other elements are not stable.
- They do not have a stable electron arrangement.
Group 1, 2, and 3 (metal atoms) will lose their valence electrons to achieve a stable electron arrangement.
- A positively charged ion is formed.
Group 5, 6, and 7 (non-metal atoms) will gain electrons to achieve a stable electron arrangement.
- A negatively charged ion is formed.
An ion is formed when an atom gains or loses electrons to form particles with a positive or negative charge.

10.4 Inside Ionic Compounds

Ionic bonding
- Formed between a metal atom and a non-metal atom, producing an ionic compound.
- Oppositely charged ions are attracted to each other by a strong electrostatic force → Ionic bond.
Example: Sodium Chloride (NaCl)
- (a) Sodium atom:
  - Proton number: 11
  - Electron configuration: 2, 8, 1
  - Loses 1 electron to achieve a stable structure: 2, 8
- (b) Chlorine atom:
  - Proton number: 17
  - Electron configuration: 2, 8, 7
  - Gains 1 electron to achieve a stable structure: 2, 8, 8
- (c) Transfer an electron from sodium to chlorine, both achieve a stable arrangement.
- (d) Oppositely-charged ions, $Na^+$ and $Cl^-$ , are strongly attracted to each other by electrostatic force → Ionic bond.
  - Sodium chloride compound is formed.
Example: Magnesium Oxide (MgO)
- (a) Magnesium atom:
  - Proton number: 12
  - Electron configuration: 2, 8, 2
  - Loses 2 electrons to achieve a stable structure: 2, 8
- (b) Oxygen atom:
  - Proton number: 8
  - Electron configuration: 2, 6
  - Gains 2 electrons to achieve a stable structure: 2, 8
- (c) Transfer 2 electrons from magnesium to oxygen, both achieve a stable arrangement.
- (d) Oppositely-charged ions, $Mg^{2+}$ and $O^{2-}$ , are strongly attracted to each other by electrostatic force → Ionic bond.
  - Magnesium oxide compound is formed.
Example: Calcium Fluoride ( $CaF_2$ )
- (a) Calcium atom:
  - Proton number: 20
  - Electron configuration: 2, 8, 8, 2
  - Loses 2 electrons to achieve a stable structure: 2, 8, 8
- (b) Fluorine atom:
  - Proton number: 9
  - Electron configuration: 2, 7
  - Gains 1 electron to achieve a stable structure: 2, 8
- (c) Transfer 1 electron from calcium to each fluorine, all achieve a stable arrangement.
- (d) Oppositely-charged ions, $Ca^{2+}$ and $F^-$ , are strongly attracted to each other by electrostatic force → Ionic bond.
  - $CaF_2$ compound is formed.
Giant ionic structure
- Ionic bonding holds all the positive and negative ions in a 3-D pattern.
Physical properties of ionic compounds
- High melting point & boiling point.
  - Contain a large number of ions; a lot of energy is needed to overcome this ionic bonding.
- Brittle.
  - Dropping a crystal of an ionic compound causes it to break between rows of ions.
  - Caused by the rearrangement of ions, like charges will repel.

10.5 Covalent Bonding

An atom can achieve a stable electron arrangement/stable structure in two ways:
- Transferring electrons from metal to non-metal → Ionic Bond
- Sharing electrons between non-metals → Covalent Bond
Covalent bonds are usually formed between atoms of non-metals, which are elements from G5-G7.
- G4 can also form covalent bonds.
Each atom contributes an equal number of electrons for sharing.
The shared pair of electrons forms a covalent bond between two atoms.
A covalent bond binds atoms together to form a group called a covalent molecule. Example: $H2$ , $Cl2$
A compound formed through covalent bonding is known as a covalent compound. Example: $NH_3$
Covalent molecule
- (1) Hydrogen molecule, $H_2$
  - Hydrogen atom, H
  - Proton: 1
  - Electron arrangement: 1
  - Needs one more electron to achieve a stable structure/stable electron arrangement.
- Two hydrogen atoms will share one pair of electrons so that both achieve a stable electron arrangement.
- A covalent bond is formed between two hydrogen atoms, forming a hydrogen molecule.
Chlorine molecule, $Cl_2$
- Chlorine atom, Cl
- Proton: 17
- Electron arrangement: 2,8,7
- Needs one more electron to achieve a stable structure/stable electron arrangement.
Two chlorine atoms will share one pair of electrons so that both achieve a stable electron arrangement.
A covalent bond is formed between two chlorine atoms, forming a chlorine molecule.
Covalent compound
- (1) Ammonia, $NH_3$
  - (i) Nitrogen atom, N
    - Proton: 7
    - Electron arrangement: 2,5
    - Needs three more electrons to achieve a stable structure/stable electron arrangement.
    - Contributes 3 electrons for sharing.
  - (ii) Hydrogen atom, H
    - Proton: 1
    - Electron arrangement: 1
    - Needs one more electron to achieve a stable structure/stable electron arrangement.
    - Contributes 1 electron for sharing.
  - (iii) One nitrogen atom shares 3 pairs of electrons with 3 hydrogen atoms so all achieve a stable structure.
  - (iv) 3 covalent bonds are formed.
  - (v) Formed ammonia compound.
Double covalent bond
- (1) Carbon dioxide, $CO_2$
  - (i) Carbon atom, C
    - Proton: 6
    - Electron arrangement: 2,4
    - Needs four more electrons to achieve a stable structure/stable electron arrangement.
    - Each carbon will contribute four electrons for sharing.
  - (ii) Oxygen atom, O
    - Proton: 8
    - Electron arrangement: 2,6
    - Needs two more electrons to achieve a stable structure/stable electron arrangement.
    - Each oxygen will contribute two electrons for sharing.
  - (iii) One carbon atom shares 4 pairs of electrons with 2 oxygen atoms so all achieve a stable structure.
  - (iv) Two double covalent bonds are formed.
  - (v) Formed carbon dioxide compound.
Triple covalent bond
- (1) Nitrogen molecules, $N_2$
  - (i) Nitrogen atom, N
    - Proton: 7
    - Electron arrangement: 2,5
    - Needs three more electrons to achieve a stable structure/stable electron arrangement.
    - Each nitrogen will contribute three electrons for sharing.
Triple covalent bonds are formed between two nitrogen atoms.

10.6 Covalent Structures

Simple molecule
- Nitrogen molecules are weakly attracted to each other.
- Little energy is needed to melt nitrogen (-210°C).
- Low melting point.
- Other substances that exist as simple molecules, such as oxygen and methane, also have low melting points.
- Most are in the gas or liquid state at room temperature (23-26°C).
Giant structure
- Carbon
  - Proton: 6
  - Needs 4 electrons to achieve a stable electron arrangement.
- Carbon cannot form a quadruple covalent bond.
- Because the distance between the nuclei for two atoms is too close.
- Repulsion will happen.
- A maximum triple covalent bond will form.
Example: Diamond
- Each carbon atom makes strong covalent bonds with four other carbon atoms.
- The pattern is repeated many times, forming a giant covalent structure.
- A three-dimensional network of atoms joined by strong covalent bonds.
- High melting point (3550°C).
- Needs a lot of energy to melt the diamond.
Example: Silicon dioxide/ Silicon (IV) oxide
- Silicon
  - Proton: 14 (2,8,4)
  - Needs 4 electrons to achieve a stable structure
- Oxygen
  - Proton: 8 (2,6)
  - Needs 2 electrons to achieve a stable structure
- All atoms are bonded by covalent bonds in a giant lattice structure.
- A lot of heat energy is required to break the thousands of strong covalent bonds in a giant structure.
Properties of giant covalent structures
- High melting and boiling points.
- Exist as a solid at room temperature.
Comparison of Ionic Compounds, Simple Covalent molecules, and Giant Covalent structures
- Ionic Compound (3-D)
 - Metal/non-metal
 - Example: NaCl
 - Melting/boiling point: high
 - Reason: Strong ionic bond
 - State: Solid
- Covalent Bond (simple molecules)
 - Non-metal
 - Example: $H2$ , $N2$ , $O_2$
 - Melting/boiling point: low
 - Reason: Weak intermolecular bond
 - State: Liquid/gas
- Giant covalent (3-D, lattice)
 - Non-metal
 - Example: Diamond, Silicon dioxide
 - Melting / boiling point: high
 - Reason: Strong covalent bond
 - State: Solid

10.7 More About Structures

Non-metal and metal → Ionic Bond
Non-metal and non-metal → Covalent Bond
Metal and metal → Metallic bond
Giant metallic structure
- Example: Sodium
- Lost their outer electrons to achieve a stable electron configuration.
Metal structure is made up from:
- Positively charged ions
- Negatively charged (delocalized) electrons
Ions are in fixed positions.
Electrons move around between ions.
Electrostatic attraction between positive ions and negative electrons.
Giant metallic structure: a three-dimensional pattern of positive metal ions held together by moving electrons
Physical properties of metallic bonding
- High strength.
- High melting point.
- Moving electrons allow metals to conduct electricity.
Summarizing structure
- Ionic
  - Giant
  - Positive ions, negative ions
  - Compound
  - Solid
  - Sodium chloride
- Metallic
  - Giant
  - Positive ions & negative electrons
  - Element
  - Solid
  - Gold
- Giant Covalent
  - Giant
  - Atoms (joined with covalent bond)
  - Element and Compound
  - Solid
  - Carbon (diamond)
- Simple Covalent
  - Simple
  - Atoms (joined with covalent bond)
  - Element and Compound
  - Gas/Liquid
  - Hydrogen, carbon dioxide
Alloy
- A combination of different elements.
- A metal that is made by mixing two or more metals, or a metal and another substance
- Example: Steel (Iron + carbon)
Different sizes of atoms in alloys
It is difficult for the layers to slide over each other.
Alloys are stronger in strength than the metal from which they are made.